CHAPTER 5 STUDY GUIDEJULY 6, 2010MORE ABOUT THE PERIODIC TABLE noble gases – have completely-filled valence electron shells; chemically unreactive (inert) due to filledvalence shellsATOMIC RADII (A-GROUP ELEMENTS ONLY) atomic radii – describes the relative sizes of atoms; adding more protons left to right, so the number ofelectrons determines atomic radii; the reason atomic radii decrease across a period is due to a shieldingeffect periodic trends for atomic radii: smaller/larger    smaller1A 8AH 2A 3A 4A 5A 6A 7A HeLi Be B C N O F NeNa Mg 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B Al Si P S Cl ArK Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br KrRb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I XeCs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At RnFr Ra larger example – arrange these elements in order of increasing atomic radii: Se, S, O, Te.(going down column 6A)O < S < Se < Te example – arrange these elements in order of increasing atomic radii: P, Cl, S, Si.(going across period 3)Cl < S < P < Si example – arrange these elements in order of increasing atomic radii: Ga, F, S, As.F < S < As < GaIONIZATION ENERGY1CHAPTER 5 STUDY GUIDEJULY 6, 2010 sometimes abbreviated as I.E. first ionization energy (IE 1) – the minimum amount of energy required to remove the most loosely boundelectron from an isolated gaseous atom to form a 1+ ion when an atom loses its first electron (like Mg+) atom(g) + energy  ion+(g) + e- Na + energy  Na+ + e- where Na has 1 unpaired electron with the electron configuration [Ne]3s1 and Na+ has0 unpaired electrons with the electron configuration [Ne] Mg(g) + 738 kJ/mol  Mg+ + e- second ionization energy (IE 2) – the amount of energy required to remove the second electron from agaseous 2+ ion when an atom loses its second electron (like Mg2+) Mg, [Ne]3s2  738 kJ/mol  Mg+, [Ne]3s1  1451 kJ/mol  Mg2+, [Ne] ion+ + energy  ion2+ + e- Mg+ + 1451 kJ/mol  Mg2+ + e- IE2 > IE1 half-filled and completely-filled orbitals are more stable, so it requires more energy to remove an electron,and you destabilize the atom by adding an electron Be 2s2  takes more energy to remove this electronB 2s22p1  than to remove this electron atoms can have third, fourth, etc. ionization energies periodic trends for ionization energy: larger/smaller    larger1A 8AH 2A 3A 4A 5A 6A 7A HeLi Be B C N O F NeNa Mg 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B Al Si P S Cl ArK Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br KrRb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I XeCs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At RnFr Ra smaller ionization energy IE2 > IE1 IE1 generally increases moving from the 1A elements to the 8A elements important exceptions at Be and Mg, N and P, etc. due to filled and half-filled subshells first ionization energies of some elements: example – arrange these elements in order of increasing first ionization energies: Sr, Be, Ca, Mg.2CHAPTER 5 STUDY GUIDEJULY 6, 2010Sr < Ca < Mg < Be example – arrange these elements in order of increasing first ionization energies: Al, Cl, Na, P.Na < Al < P < Cl example – arrange these elements in order of increasing first ionization energies: B, O, Be, N.B < Be < O < NELECTRON AFFINITY electron affinity – the amount of energy absorbed when an electron is added to an isolated gaseous atom toform an ion with a 1- charge sign conventions: if electron affinity > 0 (+), then energy is absorbed if electron affinity < 0 (–), then energy is released electron affinity is essentially a measure of an atom’s ability to form negative ions atom(g) + e- + EA  ion-(g) Mg(g) + e- + 231 kJ/mol  Mg-(g)EA = +231 kJ/mol (endothermic reaction) Br(g) + e-  Br-(g) + 323 kJ/molEA = -323 kJ/mol (exothermic reaction) periodic trends for electron affinity: smaller/larger    smaller1A 8AH 2A 3A 4A 5A 6A 7A HeLi Be B C N O F NeNa Mg 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B Al Si P S Cl ArK Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br KrRb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I XeCs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At RnFr Ra larger exceptions for filled and half-filled subshells Be 2s2  filled EA ~ 0N 2p3  half-filled EA = 0 electron affinity values of some elements:  example – arrange these elements in order of increasing electron affinities: Al, Mg, Si, Na.3CHAPTER 5 STUDY GUIDEJULY 6, 2010Si < Na < Al < Mg(Na is bizarre in this instance; otherwise, would have been Si < Al < Na < Mg) Na atom vs. Na+ ion Na atom [Ne]3s1 ) ) )e- Na+ ion [Ne]  ) )e-IONIC RADII cations (positive ions) are always smaller than their respective neutral atoms cation (positive ion) radii decrease going from left to right across a period increasing nuclear charge attracts the electrons and decreases the ionic radius anions (negative ions) are always larger than their neutral atoms anion (negative ion) radii decrease going from left to right across a period increasing electron numbers in highly-charged ions cause the electrons to repel and increasethe ionic radius periodic trends for ionic radii: smaller/larger    smaller1A 8AH 2A 3A 4A 5A 6A 7A HeLi Be B C N O F NeNa Mg 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B Al Si P S Cl ArK Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br KrRb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I XeCs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At RnFr Ra larger except it’s more like this:(Na+) (Mg2+) (Al3+) … (N3-) (O2-) (F-) more electrons fighting for fewer protons causes repulsion 7N3-7 p+10 e- ) ) - 8O2-8 p+10 e- ) ) 2- 9F-9 p+10 e- ) ) ) 3- example – arrange these elements in order of decreasing ionic radii: Ga, K, Ca.K+ > Ca2+ > Ga3+ example – arrange these elements in order of increasing ionic radii: Cl, Se, Br, S.Cl- < S2- < Br- < Se2-Cl = period 3; S = period 3; Br = period 4; Se = period 4ELECTRONEGATIVITY electronegativity – a measure of the relative tendency of an atom to attract electrons to itself whenchemically combined with another element the higher the number, the more they are attracted to each other numbers range …